5
Chemical Bonding
– Lewis Theory
C H A P T E R
C O N T E N T S
ELECTRONIC THEORY OF
VALENCE
IONIC BOND
EXAMPLES OF IONIC
COMPOUNDS
CHARACTERISTICS OF IONIC
COMPOUNDS
COVALENT BOND
CONDITIONS FOR FORMATION
OF COVALENT BOND
EXAMPLES OF COVALENT
COMPOUNDS
CHARACTERISTICS OF
COVALENT COMPOUNDS
CO-ORDINATE COVALENT BOND
EXAMPLES OF COORDINATE
COMPOUNDS OR IONS
DIFFERENCES BETWEEN IONIC
AND COVALENT BONDS
POLAR COVALENT BONDS
HYDROGEN BONDING (H-bonding)
EXAMPLES OF HYDROGENBONDED COMPOUNDS
CHARACTERISTICS OF
HYDROGEN-BOND COMPOUNDS
EXCEPTIONS TO THE OCTET
RULE
VARIABLE VALENCE
METALLIC BONDING
GEOMETRIES OF MOLECULES
VSEPR THEORY
TERMS AND DEFINITIONS
Chemical Bond
Molecules of chemical substances are made of two or more
atoms joined together by some force, acting between them. This
force which results from the interaction between the various atoms
that go to form a stable molecule, is referred to as a Chemical
Bond.
A chemical bond is defined as a force that acts between two or
more atoms to hold them together as a stable molecule.
As we will study later, there are three different types of bonds
recognised by chemists :
(1) Ionic or Electrovalent bond
(2) Covalent bond
(3) Coordinate covalent bond
There is a fourth type of bond, namely, the metallic bond
which we will consider later in this chapter.
Definition of Valence
The term valence (or valency) is often used to state the
potential or capacity of an element to combine with other elements.
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5 PHYSICAL CHEMISTRY
At one time, it was useful to define valence of an element as : the number of hydrogen atoms or twice
the number of oxygen atoms with which that element could combine in a binary compound (containing
two different elements only).
In hydrogen chloride (HCl), one atom of chlorine is combined with one atom of hydrogen and the
valence of chlorine is 1. In magnesium oxide (MgO), since one atom of magnesium holds one atom of
oxygen, the valence of magnesium is 2.
By the above definition, we would assign a valence of 2 to sulphur in H2S, but 4 to sulphur in SO2.
Some elements have fractional valence in certain compounds, while there are elements that have
variable valencies. The concept of valence as a mere number could not explain these facts. This
concept, in fact, was very confusing and has lost all value.
As already stated, there are three different type of bonds that are known to join atoms in
molecules. Although no precise definition of valence is possible, we can say that : Valence is the
number of bonds formed by an atom in a molecule.
Valence Electrons
The electrons in the outer energy level of an atom are the ones that can take part in chemical
bonding. These electrons are, therefore, referred to as the valence electrons.
The electronic configuration of Na is 2, 8, 1 and that of Cl is 2, 8, 7. Thus sodium has one valence
electron and chlorine 7. It is important to remember that for an A group element of the periodic table
(H, O, K, F, Al etc.) the group number is equal to the number of valence electrons.
Bonding and Non-bonding Electrons
The valence electrons actually involved in bond formation are called bonding electrons. The
remaining valence electrons still available for bond formation are referred to as non-bonding electrons.
Thus :
Bonding electrons
Cl Cl
Nonbonding
electrons
Lewis Symbols of Elements
A Lewis symbol of an element consists of an element’s symbol and surrounding dots to represent
the number of valence electrons. In this notation, the symbol of an element represents the nucleus
plus the inner normally filled levels (or shells) of the atom. For illustration, the symbol Na stands for
the nucleus of sodium atom plus 2, 8 electrons in the inner two levels.
Figure 5.1
The Lewis symbol Na represents the nucleus and the electrons arranged
in the inner two levels as 2, 8, minus the valence electrons.
CHEMICAL BONDING - LEWIS THEORY
153
To represent a Lewis symbol for an element, write down the symbol of the element and surround
the symbol with a number of dots (or crosses) equal to the number of valence electrons. The position
of dots around the symbol is not really of any significance. The bonding electrons are shown at
appropriate positions, while the rest of the electrons are generally given in pairs. The Lewis symbols
for hydrogen, chlorine, oxygen and sulphur may be written as :
H
Cl
O
S
The structural formulae of compounds built by union of Lewis symbols for the component atoms,
are referred to as Electron-dot formulas, or Electron-dot structures or Lewis structures. For this
purpose, the valence electrons actually involved in bond formation may be shown by crosses (x) or
dots (.) for the sake of distinction.
Now we will proceed to discuss the common types of chemical bonds in the light of the electronic
theory of valence.
ELECTRONIC THEORY OF VALENCE
As Bohr put forward his model of the atom so electronic configuration of elements was known.
G.N. Lewis and W. Kossel, working independently, used this knowledge to explain ‘why atoms joined
to form molecules’. They visualised that noble gas atoms had a stable electronic configuration, while
atoms of all other elements has unstable or incomplete electronic configuration. In 1916, they gave the
electronic theory of valence. It states that : In chemical bond formation, atoms interact by losing,
gaining, or sharing of electrons so as to acquire a stable noble gas configuration. Each noble gas,
except helium, has a valence shell of eight electrons (Table 5.1).
TABLE 5.1. ELECTRONIC CONFIGURATION OF NOBLE GASES
Noble gas
At. No.
Electrons in principal shells
He
Ne
Ar
Kr
Xe
Rn
2
10
18
36
54
86
2
2, 8
2, 8, 8
2, 8, 18, 8
2, 8, 18, 18, 8
2, 8, 18, 32, 18, 8
While atoms of noble gases possess a stable outer shell of eight electrons or octet, atoms of
most other elements have incomplete octets. They may have less than 8 electrons or in excess.
Thus, the electronic theory or valence could well be named as the Octet theory of Valence. It may
be stated as : Atoms interact by electron-transfer or electron-sharing, so as to achieve the stable
outer shell of eight electrons.
The tendency for atoms to have eight electrons in the outer shell is also known as the Octet Rule
or the Rule of Eight. Since helium has two electrons in the outer shell, for hydrogen and lithium,
having one and three (2, 1) electrons respectively, it is the Rule of two which will apply. We will see
later in this chapter that there are quite a few exceptions to the rule of eight in covalent compounds.
IONIC BOND
This type of bond is established by transfer of an electron from one atom to another. Let us
consider a general case when an atom A has one electron in the valence shell and another atom B has
seven electrons. A has one electron in excess and B has one electron short than the stable octet.
Therefore, A transfers an electron to B and in this transaction both the atoms acquire a stable electronoctet. The resulting positive ion (cation) and negative ion (anion) are held together by electrostatic
attraction.
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5 PHYSICAL CHEMISTRY
LINUS CARL PAULING
Linus Pauling won the Nobel Prize in Chemistry in 1954 for his work on chemical bonding.
He also received the Nobel Peace Prize in 1962 for his campaign against nuclear testing.
Linus Carl Pauling (February 28, 1901 – August 19, 1994) was an American quantum chemist
and biochemist. Pauling is widely regarded as the premier chemist of the twentieth century. He
pioneered the application of quantum mechanics to chemistry and in 1954 was awarded the Nobel
Prize in chemistry for his work describing the nature of chemical bonds. He also made important
contributions to crystal and protein structure determination, and was one of the founders of
molecular biology. He came near to discovering the “double helix,” the ultrastructure of DNA,
when Watson and Crick made the discovery in 1953.
Pauling received the Nobel Peace Prize in 1962 for his campaign against above-ground nuclear
testing, and is the only person to win two unshared Nobel prizes. Later in life, he became an
advocate for greatly increased consumption of vitamin C and other nutrients. He generalized his
ideas to define orthomolecular medicine, which is still regarded as unorthodox by conventional
medicine. He popularized his concepts, analyses, research and insights in several successful but
controversial books centered around vitamin C and orthomolecular medicine.
A
B
(1 Valence
electron)
(7 Valence
electrons)
+
A
+
+
A
+
B
+
or A
+
Cation
B
Anion
Electrovalent
bond
B
+
A
B
or
+
A B
The electrostatic attraction between the cation (+) and anion (–) produced by electron-transfer
constitutes an Ionic or Electrovalent bond.
The compounds containing such a bond are referred to as Ionic or Electrovalent Compounds.
CONDITIONS FOR FORMATION OF IONIC BOND
The conditions favourable for the formation of an ionic bond are :
(1) Number of valence electrons
The atom A should possess 1, 2 or 3 valence electrons, while the atom B should have 5, 6 or 7
valence electrons. The elements of group IA, IIA and IIIA satisfy this condition for atom A and those
of groups VA, VIA, and VIIA satisfy this condition for atom B.
CHEMICAL BONDING - LEWIS THEORY
155
(2) Net lowering of Energy
To form a stable ionic compound, there must be a net lowering of the energy. In other words
energy must be released as a result of the electron transfer and formation of ionic compound by the
following steps :
(a) The removal of electron from atom A (A – e– → A+) requires input of energy, which is the
ionization energy (IE). It should be low.
(b) The addition of an electron to B (B + e– → B – ) releases energy, which is the electron affinity
of B (EA). It should be high.
(c) The electrostatic attraction between A+ and B – in the solid compound releases energy, which
is the electrical energy. It should also be high.
If the energy released in steps (b) and (c) is greater than the energy consumed in step (a), the
overall process of electron transfer and formation of ionic compound results in a net release of energy.
Therefore, ionisation of A will occur and the ionic bond will be formed. For example, in case of
formation of sodium chloride (NaCl), we have
Na
–
e–
⎯⎯
→
Na +
Cl
+
e–
⎯⎯
→
Cl –
Na
+
+
Cl
–
⎯⎯
→
+
Na Cl
–
–
119 kcal
+
85 kcal
+
187 kcal
The net energy released is 187 + 85 – 119 = 153 kcal. Since the overall process results in a lowering
of energy, the ionic bond between Na and Cl will be formed.
(3) Electronegativity difference of A and B
From the line of argument used in (2), we can say that atoms A and B if they have greatly different
electronegativities, only then they will form an ionic bond. In fact, a difference of 2 or more is
necessary for the formation of an ionic bond between atoms A and B. Thus Na has electronegativity
0.9, while Cl has 3.0. Since the difference is (3.0 – 0.9) = 2.1, Na and Cl will form an ionic bond.
FACTORS GOVERNING THE FORMATION OF IONIC BOND
(1) Ionisation Energy
The ionisation energy of the metal atom which looses electron(s) should be low so that the
formation of +vely charged ion is easier. Lower the ionisation energy greater will be the tendency
of the metal atom of change into cation and hence greater will be the ease of formation of ionic
bond. That is why alkali metals and alkaline earth metals form ionic bonds easily. Out of these
two, alkali metals form ionic bonds easily as compared to alkaline earth metals. In a group the
ionisation energy decreases as we move down the group and therefore, the tendency to form
ionic bond increases in a group downward. Due to this reason Cs is the most electropositive atom
among the alkali metals.
(2) Electron Affinity
The atom which accepts the electron and changes into anion should have high electron affinity.
Higher the electron affinity more is the energy released and stable will be the anion formed. The
elements of group VI A and VII A have, in general, higher electron affinity and have high tendency to
form ionic bonds. Out of these two, the elements of group VII A (halogens) are more prone to the
formation of ionic bond than the elements of group VI A. In moving down a group the electron
affinity decreases and, therefore, the tendency to form ionic bond also decreases.
(3) Lattice Energy
After the formation of cations and anions separately, they combine to form ionic compound.
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5 PHYSICAL CHEMISTRY
+
A
+
+
B
A
B
+
Lattice energy
In this process, energy is released. It is called Lattice Energy. It may be defined as “the amount of
energy released when one mole of an ionic compound is formed from its cations and anions.”
Greater the lattice energy, greater the strength of ionic bond. The value of lattice energy depends
upon the following two factors :
(a) Size of the ions
In order to have the greater force of attraction between the cations and anions their size should
be small as the force of attraction is inversely proportional to the square of the distance between them.
(b) Charge on Ions
Greater the charge on ions greater will be the force of attraction between them and, therefore,
greater will be the strength of the ionic bond.
Necessary for the formation of an ionic bond between atoms A and B. Thus Na has electronegativity
0.9, while Cl has 3.0. Since the difference is (3.0 – 0.9) = 2.1, Na and Cl will form an ionic bond.
SOME EXAMPLES OF IONIC COMPOUNDS
Here we will discuss the formation of Lewis formula or Electron dot formula of some binary ionic
compounds, for illustration.
Sodium Chloride, NaCl
A simple sodium chloride molecule is formed from an atom of sodium (Na) and one atom of
chlorine (Cl). Na (2, 8, 1) has one valence electron, while Cl (2, 8, 7) has seven. Na transfers its valence
electron to Cl, and both achieve stable electron octet. Thus Na gives Na+ and Cl gives Cl– ion, and the
two are joined by an ionic bond.
Na
2,8,1
+
Cl
2,8,7
Na+
2,8
+
Cl
2,8,8
or
Na+ClSodium
chloride
Ionic Compounds Exist as Crystals. The (+) and (–) ions attract each other with electrostatic
force that extends in all directions. This means that ions will be bonded to a number of oppositely
charged ions around them. Therefore in solid state, single ionic molecules do not exist as such.
Rather many (+) and (–) ions are arranged systematically in an alternating cation-anion pattern called
the crystal lattice. The crystal lattice of NaCl is shown in Fig. 5.2. It will be noticed that here a large
number of Na+ and Cl– ions are arranged in an orderly fashion so as to form a cubic crystal. Each Na+
ion is surrounded by 6 Cl– ions and each Cl– ion is surrounded by 6 Na+ ions. This makes a network
of Na+ and Cl– ions which are tightly held together by electrostatic forces between them.
CHEMICAL BONDING - LEWIS THEORY
157
Although discrete molecules Na+Cl– do not exist in the solid form of ionic compounds, independent
molecules do exist in the vapour form of such compounds.
Figure 5.2
Ionic crystal of Sodium Chloride.
Magnesium Chloride, Mg 2+ Cl12- (MgCl2)
Magnesium (Mg) has two valence electrons, while chlorine (Cl) has seven. The magnesium
atom transfers its two electrons, one to each chlorine atom, and thus all the three atoms achieve the
stable octet. In this way Mg atom gives Mg2+ ion and the two Cl atoms give 2Cl1–, forming Mg 2 + Cl1–
2
(or MgCl2).
Mg
Cl
2, 8, 2
or
2, 8, 7
Mg
Mg
Cl
2, 8, 7
2Cl
2, 8
Mg
2+
Cl
2+
2, 8, 8
2Cl or
Cl
2, 8, 8
2+
Mg Cl 2
Calcium Oxide, Ca2+O2– (CaO)
Calcium (Ca) has two valence electrons, while oxygen (O) has six. Calcium atom transfers its two
valence electrons to the same oxygen atom. Thus both Ca and O achieve the stable electron-octet,
forming Ca2+ and O2– ions. Thus is obtained the molecule of calcium oxide, Ca2+O2–.
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5 PHYSICAL CHEMISTRY
xx
2+
Ca
O
Ca
2, 8, 8, 2
2, 6
2, 8, 8
2
+
O
2+
2
Ca O
or
Calcium
oxide
2, 8
Aluminium Oxide, Al32 + O 32 – (Al2O3)
Here the aluminium atom (Al) has three electrons in the valence shell (2, 8, 3), while oxygen has six
(2, 6). Two atoms of aluminium transfer their six electrons to three oxygen atoms. Thus are the electronoctets of the two Al atoms and three O atoms achieved. The two Al atoms deprived of three electrons
each, give 2Al3+ ions, while the three O atoms having gained two electrons each give 3O2– ions. In this
way, we get Al 32 + O 32 – or Al2O3.
CHARACTERISTICS OF IONIC COMPOUNDS
The ionic compounds are made of (+) and (–) ions held by electrostatic forces in a crystal lattice.
Each ion is surrounded by the opposite ions in alternate positions in a definite order in all directions.
This explains the common properties of ionic compounds.
2
O
Al
Al
+
Al
2Al
3O
2Al
3+
x
x
O
x
x
O
2
3+
O
or
O
3+
O
Al
x
x
3O
2
or
2
3+
2
2Al2 O3
Aluminium
oxide
(1) Solids at Room Temperature
On account of strong electrostatic forces between the opposite ions, these ions are locked in
their allotted positions in the crystal lattice. Since they lack the freedom of movement characteristic of
the liquid state, they are solids at room temperature.
(2) High Melting Points
Ionic compounds have high melting points (or boiling points). Since the (+) and (–) ions are
tightly held in their positions in the lattice, only at high temperature do the ions acquire sufficient
kinetic energy to overcome their attractive forces and attain the freedom of movement as in a liquid.
Thus ionic compounds need heating to high temperatures before melting.
(3) Hard and brittle
The crystals of ionic substances are hard and brittle. Their hardness is due to the strong
electrostatic forces which hold each ion in its allotted position.
These crystals are made of layers of (+) and (–) ions in alternate positions so that the opposite
ions in the various parallel layers lie over each other. When external force is applied to a layer of ions
(Fig. 5.3), with respect to the next, even a slight shift brings the like ions in front of each other. The (+)
and (–) ions in the two layers thus repel each other and fall apart. The crystal cleaves here.
CHEMICAL BONDING - LEWIS THEORY
159
Figure 5.3
(a) Two layers of (+) and (–) ions in a crystal. (b) When force is applied to one layer
it slips over the other so that similar ions come above one another and electrical
repulsions between them cause cleavage of the crystal.
(4) Soluble in water
When a crystal of an ionic substance is placed in water, the polar water molecules detach the (+)
and (–) ions from the crystal lattice by their electrostatic pull. These ions then get surrounded by
water molecules and can lead an independent existence and are thus dissolved in water. By the same
reason, non-polar solvents like benzene (C 6H 6) and hexane (C 6H 14 ) will not dissolve ionic
compounds.
Figure 5.4
Solvation of NaCl in water.
(5) Conductors of electricity
Solid ionic compounds are poor conductors of electricity because the ions are fixed rigidly in
their positions. In the molten state and in water solutions, ions are rendered free to move about. Thus
molten ionic compounds or their aqueous solutions conduct a current when placed in an electrolytic
cell.
(6) Do not exhibit isomerism
The ionic bond involving electrostatic lines of force between opposite ions, is non-rigid and nondirectional. The ionic compounds, therefore, are incapable of exhibiting stereoisomerism.
(7) Ionic reactions are fast
Ionic compounds give reactions between ions and these are very fast.
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5 PHYSICAL CHEMISTRY
COVALENT BOND
The electron transfer theory could not explain the bonding in molecules such as H2, O2, Cl2 etc.,
and in organic molecules, that had no ions. It was G.N. Lewis who suggested that two atoms could
achieve stable 2 or 8 electrons in the outer shell by sharing electrons between them. Let us consider
a general case where an atom A has one valence electron and another atom B has seven valence
electrons. As they approach each other, each atom contributes one electron and the resulting
electron pair fills the outer shell of both the atoms. Thus A acquires stable 2 electrons and B, 8
electrons in the outer shell.
The shared pair is indicated by a dash (–) between the two bonded atoms. A shared pair of
electrons constitutes a Covalent bond or Electron-pair bond.
In fact, the positive nuclei of atoms A and B are pulled towards each other by the attraction of the
shared electron pair. At the same time, the nuclei of two atoms also repel each other as do the two
electrons. It is the net attractive force between the shared electrons and the nuclei that holds the
atoms together. Thus an alternative definition of a covalent bond would be :
The attractive force between atoms created by sharing of an electron-pair.
The compounds containing a covalent bond are called covalent compounds.
CONDITIONS FOR FORMATION OF COVALENT BOND
The conditions favourable for the formation of covalent bonds are :
(1) Number of valence electrons
Each of the atoms A and B should have 5, 6 or 7 valence electrons so that both achieve the stable
octet by sharing 3, 2 or 1 electron-pair. H has one electron in the valence shell and attains duplet. The
non-metals of groups VA, VIA and VIIA respectively satisfy this condition.
(2) Equal electronegativity
The atom A will not transfer electrons to B if both have equal electronegativity, and hence
electron sharing will take place. This can be strictly possible only if both the atoms are of the same
element.
(3) Equal sharing of electrons
The atoms A and B should have equal (or nearly equal) electron affinity so that they attract the
bonding electron pair equally. Thus equal sharing of electrons will form a nonpolar covalent bond. Of
course, precisely equal sharing of electrons will not ordinarily occur except when atoms A and B are
atoms of the same element, for no two elements have exactly the same electron affinity.
SOME EXAMPLES OF COVALENT COMPOUNDS
The construction of Lewis structures of simple covalent compounds will be discussed.
Hydrogen, H2
Hydrogen molecule is made of two H atoms, each having one valence electron. Each contributes
an electron to the shared pair and both atoms acquire stable helium configuration. Thus stable H2
molecule results.
CHEMICAL BONDING - LEWIS THEORY
161
Chlorine, Cl2
Each Cl atom (2, 8, 7) has seven valence electrons. The two Cl atoms achieve a stable electron
octet by sharing a pair of electrons.
Water, H2O
Oxygen atom (2, 6) has six valence electrons and can achieve the stable octet by sharing two
electrons, one with each H atom. Thus Lewis structure of water can be written as :
Unshared
electron pair
Sharing electrons
xx
xx
H
x
Ox
x x
or
H xOx H
H
xx
xx
H
O
H
x x
Water
Ammonia, NH3
Nitrogen atom (2, 5) has five valence electrons and can achieve the octet by sharing three
electrons, one each with three H atoms. This gives the following Lewis structure for ammonia :
Sharing electrons
Unshared
electron pair
xx
Nx
x
x
x
x x
x
xx
H
or
H N H
H
N
H
x
H
H
H
Ammonia
Methane, CH4
Carbon atom (2, 4) has four electrons in the valence shell. It can achieve the stable octet by
sharing these electrons with four H atoms, one with each H atom. Thus the Lewis structure of methane
can be written as :
Sharing
electrons
C
C
C
Methane
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5 PHYSICAL CHEMISTRY
EXAMPLES OF MULTIPLE COVALENT COMPOUNDS
In many molecules, we find that in order to satisfy the octet, it becomes necessary for two atoms
to share two or three pairs of electrons between the same two atoms. The sharing of two pairs of
electrons is known as a Double bond and the sharing of three pairs of electrons a Triple bond. Let us
consider some examples of compounds containing these multiple covalent bonds in their molecules.
Oxygen, O2
The conventional Lewis structure of oxygen is written by sharing of two pairs of electrons
between two O atoms (2, 6). In this way both the O atoms achieve the octet.
Double bond
Sharing electrons
x
x
xx
O xx
x
x
O
2, 8, 6
2, 8, 6
xx
O xx O
x
x
or
xx
O
O
Oxygen
2, 8, 8 2, 8, 8
The above structure of oxygen implies that all the electrons in oxygen, O2, are paired whereby the
molecule should be diamagnetic. However, experiment shows that O2 is paramagnetic with two unpaired
electrons. This could be explained by the structure.
O
O
Although writing Lewis structures work very well in explaining the bonding in most simple
molecules, it should be kept in mind that it is simply the representation of a theory. In this case, the
theory just doesn’t work.
Nitrogen, N2
The two atoms of nitrogen (2, 5), each having five electrons in the valence shell, achieve the octet
by sharing three electron pairs between them.
+
N
N
N
or
N
N
N
Nitrogen
(3 Electrons shared
by each N atom)
Carbon Dioxide, CO2
Carbon (2, 4) has four valence electrons. It shares two electrons with each O atom (having six
valence electrons). Thus the C atom and both the O atoms achieve their octet.
O
+
x
x
C xx
+
O
O
x
x
C xx O
or
O
C
O
Carbon dioxide
CHARACTERISTICS OF COVALENT COMPOUNDS
While the atoms in a covalent molecule are firmly held by the shared electron pair, the individual
molecules are attracted to each other by weak van der Waals forces. Thus the molecules can be
separated easily as not much energy is required to overcome the intermolecular attractions. This
explains the general properties of covalent compounds.
(1) Gases, liquids or solids at room temperature
The covalent compounds are often gases, liquids or relatively soft solids under ordinary conditions.
This is so because of the weak intermolecular forces between the molecules.
(2) Low melting points and boiling points
Covalent compounds have generally low melting points (or boiling points). The molecules are
CHEMICAL BONDING - LEWIS THEORY
163
held together in the solid crystal lattice by weak forces. On application of heat, the molecules are
readily pulled out and these then acquire kinetic energy for free movement as in a liquid. For the same
reason, the liquid molecules are easily obtained in the gaseous form which explains low boiling points
of covalent liquids.
(3) Neither hard nor brittle
While the ionic compounds are hard and brittle, covalent compounds are neither hard nor brittle.
There are weak forces holding the molecules in the solid crystal lattice. A molecular layer in the crystal
easily slips relative to other adjacent layers and there are no ‘forces of repulsion’ like those in ionic
compounds. Thus the crystals are easily broken and there is no sharp cleavage between the layers on
application of external force.
(4) Soluble in organic solvents
In general, covalent compounds dissolve readily in nonpolar organic solvents (benzene, ether).
The kinetic energy of the solvent molecules easily overcomes the weak intermolecular forces.
Covalent compounds are insoluble in water. Some of them (alcohols, amines) dissolve in water
due to hydrogen-bonding.
(5) Non-conductors of electricity
Since there are no (+) or (–) ions in covalent molecules, the covalent compounds in the molten or
solution form are incapable of conducting electricity.
(6) Exhibit Isomerism
Covalent bonds are rigid and directional, the atoms being held together by shared electron pair
and not by electrical lines of force. This affords opportunity for various spatial arrangements and
covalent compounds exhibit stereoisomerism.
(7) Molecular reactions
The covalent compounds give reactions where the molecule as a whole undergoes a change.
Since there are no strong electrical forces to speed up the reaction between molecules, these reactions
are slow.
CO-ORDINATE COVALENT BOND
In a normal covalent bond, each of the two bonded atoms contributes one electron to make the
shared pair. In some cases, a covalent bond is formed when both the electrons are supplied entirely by
one atom. Such a bond is called co-ordinate covalent or dative bond. It may be defined as : a covalent
bond in which both electrons of the shared pair come from one of the two atoms (or ions). The
compounds containing a coordinate bond are called coordinate compounds.
If an atom A has an unshared pair of electrons (lone pair) and another atom B is short of two
electrons than the stable number, coordinate bond is formed. A donates the lone pair to B which
accepts it. Thus both A and B achieve the stable 2 or 8 electrons, the lone pair being held in
common.
The atom A which donates the lone pair is called the donor, while B which accepts it the acceptor.
The bond thus established is indicated by an arrow pointing from A to B. Although the arrow head
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5 PHYSICAL CHEMISTRY
indicates the origin of the electrons, once the coordinate bond is formed it is in no way different from
an ordinary covalent bond.
The molecule or ion that contains the donor atom is called the ligand.
SOME EXAMPLES OF COORDINATE COMPOUNDS OR IONS
Lewis structures of some common molecules or ions containing a coordinate covalent bond are
listed below.
Ammonium ion, NH+4
In ammonia molecule, the central N atom is linked to three H atoms and yet N has an unshared
pair of electrons. The H+ ion furnished by an acid has no electron to contribute and can accept a pair
of electrons loaned by N atom. Thus, NH3 donates its unshared electrons to H+ forming ammonium
ion.
H
H
H
N
+
H
H
Hydrogen
ion
H
+
H
+
H
N
or
H
N
H
H
H
Ammonia
molecule
(N has a lone pair)
Ammonium ion
All the N–H bonds in NH +4 are identical, once the coordinate bond N→H+ is established.
Hydronium ion, H3O+
The oxygen atom in water molecule is attached to two H atoms by two covalent bonds. There are
still two unshared pairs of electrons with the O atom. The O atom donates one of these pairs of
electrons to H+ ion and the hydronium ion is thus formed.
H
H+
O
+
H
O
H
+
H3O
or
H
H
Hydronium ion
Fluoroborate ion, B F4–
It is formed when a boron trifluoride molecule (BF3) shares a pair of electrons supplied by fluoride
ion (F–).
F
F
F
B
F
F
Fluoride
ion
Boron
trifluoride
(B acts as Acceptor)
F
B
F
or
BF4
F
Fluoroborate
ion
Addition compound of NH3 with BCl3
The N atom of ammonia molecule (NH3) has lone pair while B atom in boron trichloride (BCl3) is
short of two electrons than stable octet. An addition compound is formed as the N atom donates its
lone pair to B atom of BCl3.
CHEMICAL BONDING - LEWIS THEORY
H
H
Cl
N
B
H
Cl
Cl
H
(N is Donor) (B is Acceptor)
H
Cl
N
B
H
Cl
165
Cl
Addition compound
Nitromethane, CH3NO2
The Lewis structure of nitromethane is shown below. Here the N atom has five valence electrons,
three of which are used in forming a covalent bond with C atom and two covalent bonds with O atom.
The N atom is still left with two unshared electrons which are donated to another O atom.
H
H
H
x
O
O
x
or
C N
x
H
C
N
O
H
O
H
Nitromethane
Sulphur dioxide, SO2, and Sulphur trioxide, SO3
Sulphur achieves its octet by forming two covalent bonds with one O atom, giving SO. The S
atom in SO has two lone pairs, one of which is shared with a second O atom to form sulphur dioxide,
SO2. The S atom in SO2 still has one lone pair which it donates to a third O atom forming the sulphur
trioxide (SO3) molecule.
O
x
x
x x
xx
S xx
or
O
S
O
Sulphur dioxide
O
x x
O
S
O
O
O
S
O
Sulphur trioxide
Aluminium Chloride, Al2Cl6
Aluminium atom has three valence electrons which it shares with three Cl atoms, forming three
covalent bonds. Thus the Al atom acquires six electrons in its outer shell. Now Cl atom has three lone
pairs, one of which is donated to the Al atom of another molecule AlCl3. Thus both Al atoms achieve
octet and stable Al2Cl6 results.
Cl
Cl
Al
Al
Cl
Cl
Cl
Cl
2–
Sulphate ion, SO4
Sulphur has six valence electrons (2, 8, 6) and achieves the octet by gaining two electrons from
metal atoms (say two Na atoms). The four pairs of electrons around the S atom are then donated to
four oxygen atoms each of which has six electrons. Thus the Lewis structure for SO24 – ion may be
written as :
166
5 PHYSICAL CHEMISTRY
2
2
O
O
xx
O xx S xx O
O
or
xx
S
O
O
O
S gains two electrons
from metal atoms and
completes octet
Sulphate ion
Ozone, O3
Oxygen molecule is made of two oxygen atoms joined by two covalent bonds. Each O atom in O2
has two unshared pairs of electrons. When one pair of these is donated to a third O atom which has
only six electrons, a coordinate bond is formed. Thus the Lewis structure of ozone may be represented
as :
x
x
xx
xx
O
O xx
+
O
O
Oxygen
(2 Atoms of O)
O
O
Ozone
Carbon Monoxide, CO
Carbon atom has four valence electrons while oxygen atom has six. By forming two covalent
bonds between them, O atom achieves octet but C atom has only six electrons. Therefore O donates
an unshared pair of electron to C, and a coordinate covalent bond is established between the two
atoms. Lewis structure of CO may be written as :
C
+
x
x
xx
O xx
C
x
x
xx
O xx
C
O
COMPARISON OF IONIC AND COVALENT BONDS
Ionic Bond
Covalent Bond
1. Formed by transfer of electrons from a metal
to a non-metal atom.
2. Consists of electrostatic force between (+)
and (–) ions.
3. Non-rigid and non-directional; cannot cause
isomerism.
1. Formed by sharing of electrons between
nonmetal atoms.
2. Consists of a shared pair of electrons
between atoms.
3. Rigid and directional : causes
stereoisomerism.
Properties of Compounds
1. Solids at room temperature.
2. High melting and boiling points.
3. Hard and brittle.
4. Soluble in water but insoluble in organic
solvents.
5. Conductors of electricity
6. Undergo ionic reactions which are fast.
Properties of Compounds
1. Gases, liquids or soft solids.
2. Low melting and boiling points.
3. Soft, much readily broken
4. Insoluble in water but soluble in organic
solvents.
5. Non-conductors of electricity.
6. Undergo molecular reactions which are slow.
CHEMICAL BONDING - LEWIS THEORY
167
POLAR COVALENT BONDS
In the H2 or Cl2 molecule, the two electrons constituting the covalent bond are equally shared by
the two identical nuclei. Due to even distribution of (+) and (–) charge, the two bonded atoms remain
electrically neutral. Such a bond is called nonpolar covalent bond. However, when two different atoms
are joined by a covalent bond as in HCl, the electron pair is not shared equally.
Nonpolar covalent
bond
Cl
Polar covalent
bond
Cl
H
Bonding pair ( : )
equally shared
Cl
Bonding pair ( : )
unequally shared
Due to a greater attraction of one nucleus (Cl) for the electrons, the shared pair is displaced
towards it. This makes one end of the bond partially positive (δ+ ) and the other partially negative (δ–).
H
Cl
or
H
Cl
A covalent bond in which electrons are shared unequally and the bonded atoms acquire a partial
positive and negative charge, is called a polar covalent bond.
A molecule having partial positive and negative charge separated by a distance is commonly
referred to as a Dipole (two poles). The dipole of a bond is indicated by an arrow from positive to
negative end with a crossed tail as shown above in HCl molecule.
Since two atoms of different elements do not have exactly the same attraction for electrons in a
bond, all bonds between unlike atoms are polar to some extent. The amount of polarity of a bond is
determined by the difference of electronegativity (or tendency to attract electrons) of the two bonded
atoms. The greater the difference of electronegativity between two atoms, greater the polarity. A
graph showing the % age ionic character and difference in electronegativity between the two atoms is
shown in Fig. 5.5.
As a matter of fact, if this difference is around 1.9 and 2.9, the bond is generally ionic, meaning
that one atom has gained complete control of the electron pair in the bond.
The percentage ionic character of a bond can be calculated by using the equation
%age ionic character = 16 [XA – XB] + 3.5 [XA – XB]2
This equation was given by Hannay and Smith.
5 PHYSICAL CHEMISTRY
168
100% Ionic
character
100
% Ionic character
80
60
50% Ionic
character
40
20
0
0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
Difference in electronegativity
Figure 5.5
Graph between % ionic character and
difference in electronegativity.
SOLVED PROBLEM. Calculate the percentage ionic character of C–Cl bond in CCl4 if the
electronegativities of C and Cl are 3.5 and 3.0 respectively.
SOLUTION
∴
%age ionic character = 16 [XA – XB] +3.5 [XA + XB]2
Give XA = 3.5 and XB = 3.0
%age ionic character = 16 (3.5 – 3.0) + 3.5 (3.5 – 3.0)2
= 8.0 + 0.875
= 8.875%
Examples of Polar Covalent Bonds
Water molecule (H2O) contains two O–H covalent bonds. The electronegativity of O is 3.5 and
that of H is 2.1. Thus both the bonds are polar and water has a polar molecule.
δ+
H
Polar covalent
bond
O
δ+
H
Water
H
δ+
N
δ+
δ
H
Ammonia
δ+
H
δ
In ammonia molecule, there are three N–H bonds. The electronegativity of N is 3.0 and that of H
is 2.1. Therefore all the N–H bonds are polar and ammonia has a polar molecule.
The electronegativity of fluorine (F) is 4.0 and that of H is 2.1. The difference of electronegativities
being very great, the molecule H–F has a strong dipole.
CHEMICAL BONDING - LEWIS THEORY
169
HYDROGEN BONDING (H-Bonding)
When hydrogen (H) is covalently bonded to a highly electronegative atom X (O, N, F), the shared
electron pair is pulled so close to X that a strong dipole results.
X
H
or
X
H
Dipole
Since the shared pair is removed farthest from H atom, its nucleus (the proton) is practically
exposed. The H atom at the positive end of a polar bond nearly stripped of its surrounding electrons,
exerts a strong electrostatic attraction on the lone pair of electrons around X in a nearby molecule.
Thus :
Electrostatic
attraction
X
H +
X
H
X
H
X
H
Hydrogen
bond
or
X
H
X
H
The electrostatic attraction between an H atom covalently bonded to a highly electronegative
atom X and a lone pair of electrons of X in another molecule, is called Hydrogen Bonding.
Hydrogen bond is represented by a dashed or dotted line.
POINTS TO REMEMBER
(1) Only O, N and F which have very high electronegativity and small atomic size, are
capable of forming hydrogen bonds.
(2) Hydrogen bond is longer and much weaker than a normal covalent bond. Hydrogen bond
energy is less than 10 kcal/mole, while that of covalent bond is about 120 kcal/mole.
(3) Hydrogen bonding results in long chains or clusters of a large number of ‘associated’
molecules like many tiny magnets.
(4) Like a covalent bond, hydrogen bond has a preferred bonding direction. This is attributed
to the fact that hydrogen bonding occurs through p orbitals which contain the lone pair
of electrons on X atom. This implies that the atoms X–H...X will be in a straight line.
CONDITIONS FOR HYDROGEN BONDING
The necessary conditions for the formation of hydrogen bonding are
(1) High electronegativity of atom bonded to hydrogen
The molecule must contain an atom of high electronegativity such as F, O or N bonded to
hydrogen atom by a covalent bond. The examples are HF, H2O and NH3.
(2) Small size of Electronegative atom
The electronegative atom attached to H-atom by a covalent bond should be quite small. Smaller
the size of the atom, greater will be the attraction for the bonded electron pair. In other words, the
polarity of the bond between H atom and electronegative atom should be high. This results in the
formation of stronger hydrogen bonding. For example, N and Cl both have 3.0 electronegativity. But
hydrogen bonding is effective in NH3 in comparison to that in HCl. It is due to smaller size of N atom
than Cl atom.
170
5 PHYSICAL CHEMISTRY
EXAMPLES OF HYDROGEN-BONDED COMPOUNDS
When hydrogen bonding occurs between different molecules of the same compound as in HF,
H2O and NH3, it is called Intermolecular hydrogen bonding. If the hydrogen bonding takes place
within single molecule as in 2-nitrophenol, it is referred to as Intramolecular hydrogen bonding. We
will consider examples of both types.
Hydrogen Fluoride, HF
The molecule of HF contains the strongest polar bond, the electronegativity of F being the
highest of all elements. Therefore, hydrogen fluoride crystals contain infinitely long chains of H–F
molecules in which H is covalently bonded to one F and hydrogen bonded to another F. The chains
possess a zig-zag structure which occurs through p orbitals containing the lone electron pair on F
atom.
Hydrogen bond
H
F
H
F
H
F
Hydrogen fluoride molecules
Water, H2O
In H2O molecule, two hydrogen atoms are covalently bonded to the highly electronegative O
atom. Here each H atom can hydrogen bond to the O atom of another molecule, thus forming large
chains or clusters of water molecules.
Hydrogen
bond
H
O
H
Water
molecule
H
O
H
H
O
H
H
O
H
Liquid water
Each O atom still has an unshared electron pair which leads to hydrogen bonding with other water
molecules. Thus liquid water, in fact, is made of clusters of a large number of molecules.
Ammonia, NH3
In NH3 molecules, there are three H atoms covalently bonded to the highly electronegative N
atom. Each H atom can hydrogen bond to N atom of other molecules.
CHEMICAL BONDING - LEWIS THEORY
171
Hydrogen bond
H
H
N
H
H
H
H
N
H
H
H
N
H
H
N
H
Ammonia
molecule
2-Nitrophenol
Here hydrogen bonding takes place within the molecule itself as O–H and N–H bonds are a part
of the same one molecule.
TYPES OF HYDROGEN-BONDING
Hydrogen bonding is of two types :
(1) Intermolecular Hydrogen bonding
This type of hydrogen bonding is formed between two different molecules of the same or different
substances e.g. hydrogen bonding in HF, H2O, NH3 etc. It is shown in the following diagram (Fig. 5.6).
Hydrogen bond
H
F
H
F
H
F
Hydrogen fluoride molecule
Hydrogen bond
Hydrogen
bond
O
H
H
O
H
H
H
Water molecule
O
H
H
H
H
N
H
N
H
Ammonia molecule
Figure 5.6
Intermolecular hydrogen bonding in HF, H2 O and NH3 .
H
H
N
H
172
5 PHYSICAL CHEMISTRY
This type of hydrogen bonding results in the formation of associated molecules. Generally
speaking, the substances with intermolecular hydrogen bonding have high melting points, boiling
points, viscosity, surface tension etc.
(2) Intramolecular Hydrogen bonding
This type of hydrogen bonding is formed between the hydrogen atom and the electronegative atom
present within the same molecule. It results in the cyclisation of the molecule. Molecules exist as discrete
units and not in associated form. Hence intramolecular hydrogen bonding has no effect on physical
properties like melting point, boiling point, viscosity, surface tension, solubility etc. For example
intramolecular hydrogen bonding exists in o-nitrophenol, 2-nitrobenzoic acid etc. as shown below :
Figure 5.7
Intramolecular hydrogen bonding.
CHARACTERISTICS OF HYDROGEN-BONDED COMPOUNDS
(1) Abnormally high boiling and melting points
The compounds in which molecules are joined to one another by hydrogen bonds, have unusually
high boiling and melting points. This is because here relatively more energy is required to separate
the molecules as they enter the gaseous state or the liquid state. Thus the hydrides of fluorine (HF),
oxygen (H2O) and nitrogen (NH3) have abnormally high boiling and melting points compared to
other hydrides of the same group which form no hydrogen bonds. In Fig. 5.8 are shown the boiling
points and melting points of the hydrides of VIA group elements plotted against molecular weights.
It will be noticed that there is a trend of decrease of boiling and melting points with decrease of
molecular weight from H2Te to H2S. But there is a sharp increase in case of water (H2O), although it has
the smallest molecular weight. The reason is that the molecules of water are ‘associated’ by hydrogen
bonds between them, while H2Te, H2Se and H2S exist as single molecules since they are incapable of
forming hydrogen bonds.
CHEMICAL BONDING - LEWIS THEORY
H2 O
o
100
Temperature oC
173
Melting
points
Boiling
points
H2 O
o
H2 Te
H2 Se
H2 S
0
H2 Te
H2 Se
H2 S
o
–100
0
60
120
Molecular weights
Figure 5.8
Boiling and melting point curves of the hydrides of VIA group showing abrupt
increase for water (H2O) although it has the lowest molecular weight.
(2) High solubilities of some covalent compounds
The unexpectedly high solubilities of some compounds containing O, N and F, such as NH3 and
CH3OH in certain hydrogen containing solvents are due to hydrogen bonding. For example, ammonia
(NH3) and methanol (CH3OH) are highly soluble in water as they form hydrogen bonds.
H
H
N
H
Ammonia
Hydrogen bond
H
O
H
Water
H
H
Hydrogen bond
C
O
H
H
Methanol
H
O
H
Water
(3) Three dimensional crystal lattice
As already stated, hydrogen bonds are directional and pretty strong to form three dimensional
crystal lattice. For example, in an ice crystal the water molecules (H2O) are held together in a tetrahedral
network and have the same crystal lattice as of diamond. This is so because the O atom in water has
two covalent bonds and can form two hydrogen bonds. These are distributed in space like the four
covalent bonds of carbon. The tetrahedral structural units are linked to other units through hydrogen
bonds as shown in Fig. 5.6.
Since there is enough empty space in its open lattice structure ice is lighter than water, while most
other solids are heavier than the liquid form.
Water as an Interesting Liquid
Water is very interesting solvent with unusual properties. It dissolves many ionic compounds
and polar organic compounds. It has high heat of vaporisation, high heat of fusion, high specific heat
with melting point 273 K and boiling point 373 K. Its structure as shown above is very interesting as
it explains many properties :
174
5 PHYSICAL CHEMISTRY
(1) Ice (solid) is lighter than water (Liquid)
The structure of water is tetrahedral in nature. Each oxygen atom is linked to two H-atoms by
covalent bonds and other two H-atoms by hydrogen bonding. In this solid state (Ice), this tetrahedral
structure is packed resulting in open cage like structure with a number of vacant space. Hence in this
structure the volume increases for a given mass of liquid water resulting in lesser density. Due to this
reason ice floats on water.
(2) Maximum density of water at 277 K (4ºC)
On melting ice, the hydrogen bonds break and water molecules occupy the vacant spaces. This
results in decrease in volume and increase in density (d = m/v). Hence density of water keeps on
increasing when water is heated. This continues upto 277 K (4ºC). Above this temperature water
molecules start moving away from one another due to increase in kinetic energy. Due to this volume
increases again and density starts decreasing. This behaviour of water is shown in the fig. 5.9.
Density
1.0
273
274
275
276
277
278
279
280
281
282
Temperature (K)
Figure 5.9
A plot of density versus temperature (water).
EXCEPTIONS TO THE OCTET RULE
For a time it was believed that all compounds obeyed the Octet rule or the Rule of eight. However,
it gradually became apparent that quite a few molecules had non-octet structures. Atoms in these
molecules could have number of electrons in the valence shell short of the octet or in excess of the
octet. Some important examples are :
(1) Four or six electrons around the central atom
A stable molecule as of beryllium chloride, BeCl2, contains an atom with four electrons in its outer
shell.
x
Be x
+
2
Cl
Cl
Be
Cl
(4 Electrons about Be)
The compound boron trifluoride, BF3, has the Lewis structure :
Cl
x
x
B
x
+
3 Cl
Cl
B
Cl
(6 Electrons about B)
CHEMICAL BONDING - LEWIS THEORY
175
The boron atom has only six electrons in its outer shell.
Beryllium chloride and boron trifluoride are referred to as electron-deficient compounds.
(2) Seven electrons around the central atom
There are a number of relatively stable compounds in which the central atom has seven electrons
in the valence shell. A simple example is chlorine dioxide, ClO2.
O
+
Cl
+
O
O
Cl
O
Chlorine dioxide
The chlorine atom in ClO2 has seven electrons in its outer shell.
Methyl radical (CH3) has an odd electron and is very short lived. When two methyl free radicals
collide, they form an ethane molecule (C2H6) to satisfy the octet of each carbon atom. Any species with
an unpaired electron is called a free radical.
H
2H
H
C
H
(C has 7 electrons)
H
H
C
C
H
H
H
Ethane
(3) Ten or more electrons around the central atom
Non-metallic elements of the third and higher periods can react with electronegative elements to
form structures in which the central atom has 10, 12 or even more electrons. The typical examples are
PCl5 and SF6.
Cl
F
F
Cl
F
Cl
S
P
F
Cl
Cl
(10 Electrons about P)
F
F
(12 Electrons about S)
The molecules with more than an octet of electrons are called superoctet structures.
In elements C, N, O and F the octet rule is strictly obeyed because only four orbitals are available
(one 2s and three 2p) for bonding. In the elements P and S, however, 3s, 3p, and 3d orbitals of their
atoms may be involved in the covalent bonds they form. Whenever an atom in a molecule has more
than eight electrons in its valence shell, it is said to have an expanded octet.
VARIABLE VALENCE
Some elements can display two or more valences in their compounds. The transition metals
belong to this class of elements. The Electronic Structure of some of these metals is given below :
176
5 PHYSICAL CHEMISTRY
TABLE 5.2. ELECTRONIC STRUCTURE OF THE TWO OUTERMOST SHELLS OF
SOME TRANSITION METALS
Sc
Cr
Mn
Fe
Co
Cu
3d 1 4s 2
3d 5 4s 1
3d 5 4s 2
3d 6 4s 2
3d 7 4s 2
3d 10 4s 1
Ag
La
Os
Ir
Pt
Au
4d 10 5s 1
5d 1 6s 2
5d 6 6s 2
5d 7 6s 2
5 d 96 s 2
5d 10 6s 1
Most of the transition metals have one or two outer-shell electrons and they form monovalent or
bivalent positive ions. But because some of the d electrons are close in energy to the
outermost electrons, these can also participate in chemical bond formation. Thus transition metals
can form ions with variable valence. For example, copper can form Cu1+ and Cu2+ ions and iron can
form Fe2+ and Fe3+ ions.
The complete electronic configuration of an iron atom is
Fe = 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
It can form Fe2+ by losing two 4s electrons,
Fe2+ = 1s2 2s2 2p6 3s2 3p6 3d 6
When iron loses two 4s electrons and one of the three 3d electrons, if forms Fe3+ ion
Fe3+ = 1s2 2s2 2p6 3s2 3p6 3d 5
1+
2+
Copper form Cu and Cu ions by losing one 4s electron, and one 4s and 3d electron respectively
Cu = 1s2 2s2 2p6 3s2 3p6 3d10 4s1
Cu1+ = 1s2 2s2 2p6 3s2 3p6 3d10
Cu2+ = 1s2 2s2 2p6 3s2 3p6 3d 9
It may be noted that the structures of Fe2+, Fe3+, Cu1+, Cu2+, Cr3+, etc., are not isoelectronic with
any of the noble gases, and hence the d electrons being unstable are available for bond formation.
(The atoms and ions that have the same number of electrons are said to be Isoelectronic).
METALLIC BONDING
The valence bonds that hold the atoms in a metal crystal together are not ionic, nor are they
simply covalent in nature. Ionic bonding is obviously impossible here since all the atoms would tend
to give electrons but none are willing to accept them. Ordinary covalent bonding is also ruled out as,
for example, sodium atom with only one outer-shell electron could not be expected to form covalent
bonds with 8 nearest neighbouring atoms in its crystal. The peculiar type of bonding which holds the
atoms together in metal crystal is called the Metallic Bonding.
Many theories have been proposed to explain the metallic bonding. Here we will discuss the
simplest of these : The Electron Sea Model.
THE ELECTRON SEA MODEL
Metal atoms are characterised by :
(1) Low ionization energies which imply that the valence electrons in metal atoms can easily be
separated.
(2) A number of vacant electron orbitals in their outermost shell. For example, the magnesium atom
with the electron configuration 1s2 2s2 2p6 3s2 3p0 has three vacant 3p orbitals in its outer electron
shell.
CHEMICAL BONDING - LEWIS THEORY
177
There is considerable overlapping of vacant orbitals on one atom with similar orbitals of adjacent
atoms, throughout the metal crystal. Thus it is possible for an electron to be delocalized and move
freely in the vacant molecular orbital encompassing the entire metal crystal. The delocalized electrons
no longer belong to individual metal atoms but rather to the crystal as a whole.
As a result of the delocalization of valence electrons, the positive metal ions that are produced,
remain fixed in the crystal lattice while the delocalized electrons are free to move about in the vacant
space in between. The metal is thus pictured as a network or lattice of positive ions of the metal
immersed in a ‘sea of electrons’ or ‘gas of electrons’. This relatively simple model of metallic bonding
is referred to as the Electron Sea model or the Electron Gas model (Fig. 5.10.)
Figure 5.10
The Electron Sea model of metallic bonding.
A metallic bond is the electrostatic force of attraction that the neighbour positive metallic ions
have for the delocalized electrons.
+
ee-
+
The electron sea model of metallic bonding explains fairly well the most characteristic physical
properties of metals.
(1) Luster or Reflectivity. The delocalized mobile electrons of the ‘electron sea’ account for this
property. Light energy is absorbed by these electrons which jump into higher energy levels and return
immediately to the ground level. In doing so, the electrons emit electromagnetic radiation (light) of the
same frequency. Since the radiated energy is of same frequency as the incident light, we see it as a
reflection of the original light.
(2) Electric Conductivity. Another characteristic of metals is that they are good conductors of
electricity. According to the electron sea model, the mobile electrons are free to move through the
vacant space between metal ions. When electric voltage is applied at the two ends of a metal wire, it
causes the electrons to be displaced in a given direction. The best conductors are the metals which
attract their outer electrons the least (low ionization energy) and thus allow them the greatest freedom
of movement.
178
5 PHYSICAL CHEMISTRY
e-
e-
ee- Exiting
electricity
eEntering eelectricity
e-
e-
e-
e-
e-
Figure 5.11
Electrical conductivity by flow of electrons
based on Electron Sea model.
(3) Heat Conductivity. If a metal is heated at one end, the heat is carried to the other end. The
mobile electrons in the area of the ‘electron sea’ around one end of the metal easily absorb heat
energy and increase their vibrational motion. They collide with adjacent electrons and transfer the
added energy to them. Thus the mobility of the electrons allows heat transfer to the other end
(Fig. 5.12).
e-
e-
e-
e-
e-
Heat
energy
e-
e-
e-
e-
Figure 5.12
Heat conduction through a metal.
(4) Ductility and Malleability. The ductility and malleability of metals can also be explained by
the electron sea model. In metals the positive ions are surrounded by the sea of electrons that ‘flows’
around them. If one layer of metal ions is forced across another, say by hammering, the internal
structure remains essentially unchanged (Fig. 5.13). The sea of electrons adjusts positions rapidly
and the crystal lattice is restored. This allows metals to be ductile and malleable. However, in ionic
crystals of salts e.g., sodium chloride, displacement of one layer of ions with respect to another brings
like charged ions near to each other. The strong repulsive forces set up between them can cause the
ionic crystals to cleave or shatter. Thus ionic crystals are brittle.
ee-
e-
e-
e-
e-
e-
e-
e-
Before
e-
e-
e-
e-
eAfter
Figure 5.13
When force is applied to the upper layer of cations it slips
to the right without changing the environments.
illustration of malleability and ductility.
e-
e-
CHEMICAL BONDING - LEWIS THEORY
179
(5) Electron Emission. When enough heat energy is applied to a metal to overcome the attraction
between the positive metal ions and an outer electron, the electron is emitted from the metallic atom.
When the frequency and, therefore, the energy of the light that strikes the metal is great enough to
overcome the attractive forces, the electron escapes from the metal with a resultant decrease in the
energy of the incident photon (Photoelectric effect).
GEOMETRIES OF MOLECULES
So far we have depicted molecules by Lewis structures in the flat plane of paper. But all
molecules containing three or more atoms are three-dimensional. The shape of a particular molecule
is determined by the specific arrangement of atoms in it and the bond angles. Molecular shapes may
be linear, bent (or angular), trigonal planar, pyramidal or tetrahedral.
The shapes of molecules can be determined in the laboratory by modern methods such as X-ray
and electron diffraction techniques. Molecular shapes are important because they are helpful in the
investigation of molecular polarity, molecular symmetry or asymmetry. Physical and chemical properties
of compounds depend on these factors. VSEPR theory throws light on the three dimensional shapes
of molecules.
VSEPR THEORY
The Lewis structure of a molecule tells us the number of pairs of electrons in the valence shell of
the central atom. These electron pairs are subject to electrostatic attractions between them. On this
basis, R.G.Gillespie (1970) proposed a theory called the Valence-Shell Electron Pair Repulsion or
VSEPR (pronounced as ‘Vesper’) theory. It states that : The electron pairs (both lone pairs and
shared pairs, surrounding the central atom will be arranged in space as far apart as possible to
minimise the electrostatic repulsion between them.
90
o
120
o
180
o
Central
atom
Two electron
o
pairs at 90
Two electron
o
pairs at 120
Two electron
o
pairs at 180
Figure 5.14
o
o
Arrangement of two electron pairs on circle at 90 , 120o and at 180 .
o
Placement of electron pairs at 180 puts them the farthest apart,
thereby minimising the electrostatic repulsion.
Let us consider the simplest case of an atom with two electron pairs. We wish to place the
electron pairs on the surface of a sphere such that they will be as far apart as possible so as to
minimise repulsion between them. Fig. 5.14 illustrates it by showing some possible placements of the
two electron pairs. The arrangement in which the electron pair-central atom-electron pair angles is
180º, makes the electron pairs farthest apart. This arrangement is called linear because the electron
pairs and the central atom are in a straight line.
VSEPR theory is simple but remarkably powerful model for predicting molecular geometries and
bond angles. While working out the shapes of molecules from this theory, it must be remembered :
(1) Multiple bonds behave as a single electron-pair bond for the purpose of VSEPR. They
represent a single group of electrons.
5 PHYSICAL CHEMISTRY
180
(2) Order of repulsions between lone pair and lone pair (lp-lp), lone pair and bonding pair (lpbp), and bonding pair and bonding pair (bp-bp) is
lp – lp > > lp – bp > bp – bp
When a molecule has lone pairs of electrons, the bonding electron pairs are pushed closer and
thus the bond angle is decreased.
Now we proceed to work out the shapes of some common molecules with the help of VSEPR
theory.
(1) Linear Molecules
(a) Beryllium chloride, BeCl2. It has the Lewis structure
Repel one
another
Cl
Be
180
Cl
Cl
Lewis structure
o
Be
Cl
VSEPR Model
Figure 5.15
Geometry of BeCl2 molecule.
The central atom Be has two bonding electron pairs and no unshared electron. According to
VSEPR theory, the bonding pairs will occupy positions on opposite sides of Be forming an angle of
180º. An angle of 180º gives a straight line. Therefore, BeCl2 molecule is linear. In general, all molecules
as A–B–A which have only two bonds and no unshared electrons are linear.
(b) Carbon dioxide, CO2. It has the structure
Counts the same
as single bond
O
C
O
Lewis structure
O
180
C
Modified
structure
O
O
C
o
O
VSEPR Model
Figure 5.16
Geometry of CO2 molecule.
The central C atom has no unshared electron. We know that a double bond counts the same as a
single bond in VSEPR model. Thus CO2 is a linear molecule.
Similarly, it can be shown that hydrogen cyanide (H – C ≡ N) and acetylene (H – C ≡ C – H) are
linear molecules.
(2) Trigonal Planar Molecules
(a) Boron trifluoride, BF3. Its Lewis structure shown that the central atom B has three bonding
electron pairs and no unshared electrons. VSEPR theory says that the three bonding electron pairs
will be as far apart as possible. This can be so if these electron pairs are directed to the corners of an
equilateral triangle. Thus VSEPR model of BF3 molecule has three F atoms at the corners of the triangle
with B atom at its centre. All the four atoms (three F and one B) lie in the same plane. Therefore, the
shape of such a molecule is called trigonal planar. The bond angle is 120º.
CHEMICAL BONDING - LEWIS THEORY
F
181
F
Equilateral
triangle
F
F
B
o
or
120
B
B
F
Lewis formula
F
F
F
F
VSEPR Model
Figure 5.17
Geometry of BF3 molecule.
(b) Sulphur trioxide, SO3. In the Lewis structure of SO3, the central S atom is joined with two O
atoms by covalent bonds. The third O atom is joined with S by a double bond. But a double bond is
counted as a single electron pair for the purpose of VSEPR model. Therefore, in effect, S has three
electron pairs around it. Thus like BF3, SO3 has trigonal planar geometry.
Counted as
single electron pair
Surrounded by
3 electron pairs
O
O
O
O
O
S
S
S
O
O
Lewis structure
Modified
structure
120 o
O
O
VSEPR Model
(trigonal, planar)
Figure 5.18
Geometry of SO3 molecule.
(3) Tetrahedral Molecules
(a) Methane, CH4. Lewis structure of methane shows that the central C atom has four bonding
electron pairs. These electron pairs repel each other and are thus directed to the four corners of a
regular tetrahedron. A regular tetrahedron is a solid figure with four faces which are equilateral triangles.
All bond angles are 109.5º.
Tetrahedron
H
H
C
H
Has four
electron pairs
H
C
C
109.5
o
H
H
H
H
Lewis structure
Electron pairs directed
to the corners of a
tetrahedron
VSEPR Model
(tetrahedral)
Figure 5.19
Geometry of CH4 molecule.
Similarly, CCl4 in which the central C atom is bonded to four other atoms by covalent bonds has
tetrahedral shape.
182
5 PHYSICAL CHEMISTRY
(b) Ammonium ion, NH4+ , and Sulphate ion, SO42– . The N atom in NH +4 and S atom in SO24 –
have four electron pairs in the valence shell. These are directed to the corners of a tetrahedron for
maximum separation from each other. Thus both NH +4 and SO 24 – have tetrahedral shape.
H
109.5
O
o
N
109.5
+1
o
S
H
H
O
O
O
H
VSEPR Model of
+2
+
NH 4
2
VSEPR Model of SO 4
Figure 5.20
2
+
Geometry of NH 4 ion and SO4 ion.
(4) Pyramidal Molecules
(a) Ammonia molecule. The Lewis structure of NH3 shows that the central N atom has three
bonding electrons and one lone electron pair. The VSEPR theory says that these electron pairs are
directed to the corners of a tetrahedron. Thus we predict that H–N–H bond angle should be 109.5º.
But the shape of a molecule is determined by the arrangement of atoms and not the unshared electrons.
Thus, if we see only at the atoms, we can visualise NH3 molecule as a pyramid with the N atom located
at the apex and H atoms at the three corners of the triangular base.
According to VSEPR theory, a lone pair exerts greater repulsion on the bonding electron pairs
than the bonding pairs do on each other. As a result, the bonds of NH3 molecule are pushed slightly
closer. This explains why the observed bond angle H–N–H is found to be 107.3º instead of 109.5º
predicted from tetrahedral geometry.
Lone pair
Pyramid
H
N
H
N
H
N
H
H
H
o
107.3 H
Lewis formula
H
Pyramidal shape
of NH3
Figure 5.21
Geometry of NH3 molecule.
H
Pyramidal shape with
bond angle decreased
to 107.3 o
All molecules in which the N atom is joined to three other atoms by covalent bonds, have
pyramidal shape. For example, amines RNH2, R2NH and R3N have pyramidal shape.
CHEMICAL BONDING - LEWIS THEORY
183
More
repulsion
N
N
o
107.3
109.5
o
Less
repulsion
Tetrahedral geometry
predicts H—N—H
o
angle of 109.5
The bond angle decreases
because of greater repulsion
of lone electron pair
Figure 5.22
Why the angle H—N—H in NH3 molecule is 107.3o
o
while the tetrahedral angle is 109.5 ?
(b) Phosphorus trichloride, PCl3. The structural formula indicates that the central phosphorus
atom has three bonding electron pairs and one lone electron pair. Thus, like NH3 it has pyramidal
shape and the observed bond angle Cl–P–Cl is 100º.
P
P
Cl
o
Cl
0
10
Cl
Cl
Cl
Cl
Pyramidal PCl3 molecule
Lewis formula
Figure 5.23
Geometry of PCI3 molecule.
(5) Bent or Angular Molecules
(a) Water, H2O. In the structural formula of H2O, the O atom is bonded to two H atoms by
covalent bonds and has two lone pairs. Thus O is surrounded by two bonding electron pairs and two
Has 4
electron pairs
H
O
O
H
O
H
Predicted
bond angle
109.5 o
H
H
105
o
H
Bent molecule of H2 O with
o
observed bond angle 105
Figure 5.24
Geometry of H2O molecule.
unshared electron pairs. VSEPR theory says that in order to secure maximum separation between
them, the four electron pairs are directed to the corners of a tetrahedron. If we look at the three atoms
(and ignore the unshared pairs), the atoms HOH lie in the same plane and the predicted bond angle is
109.5º. But with two unshared pairs repelling the bonding pairs, the bond angle is compressed to 105º,
the experimental value. Thus the H2O molecule is flat and bent at an angle at the O atom. Such a
molecule is called a bent molecule or angular molecule.
184
5 PHYSICAL CHEMISTRY
(b) Sulphur dioxide, SO2. The Lewis structure of SO2 is given below. The S atom is bonded to
one O by a double bond and to the other O by a single bond. It has an unshared electron pair. In
VSEPR model a double bond is counted as a single electron pair. That way, the S atom is surrounded
by three electron pairs, two bonding pairs and one unshared pair. For maximum separation the three
electron pairs are directed to the corners of an equilateral triangle. The predicted bond angle is 120º.
But with the unshared electron pair repelling the bonding electron pairs, the bond angle is actually
reduced somewhat. Thus SO2 has a planar bent molecule with the observed bond angle 119.5º.
O
S
O
S
S
Lewis structure
120
O
o
VSEPR Model
O
Figure 5.25
Geometry of SO2 molecule.
O
O
Bent molecule of SO2
with bond angle 119.5 o
SUMMARY : SHAPES OF MOLECULES
Methane (CH4)
Ammonia (NH3)
Water (H2O)
Hydrogen fluoride (HF)
The directional nature of covalent bonds is shown in the diagrams of molecules above. The
shape of the methane molecule is tetrahedral because the four bonding pairs of electrons repel each
other equally, and the equilibrium position of all four bonding electron pairs is tetrahedral.
HOW TO WORK OUT THE SHAPE OF A MOLECULE
It is possible to work out the shape of a small molecule that has a formula XYn by applying a few
simple rules. We will use ammonia as an example to illustrate the idea.
Rule 1 First find the number of bonding pairs of electrons in the molecule. The number of bonding
pairs of electrons in the molecule NH3 can be seen in the formula. There must be three
bonding pairs of electrons holding the three hydrogens onto the nitrogen.
Rule 2 Find the number of valence electrons (electrons in the outer energy level) on an atom of
the central atom (The one of which there is only one.) Nitrogen is in group V, so the
nitrogen has five electrons in the outer energy level.
Rule 3 Find the number of lone pairs on the central atom by subtracting the
number of bonding pairs (3) from the valence electrons (5) to find the
number of electrons (2) that will make up lone pairs of electrons.
Divide this number by 2 to find the number of lone pairs, 2/2 = 1.
Rule 4 Distribute all the electron pairs around the central atom and learn the
angles they will make from molecules with no lone pairs.
Rule 5 Learn that the repulsion between lone pairs of electrons is greater
than the repulsion between bonding pairs, and subtract 2o from the
bond angles for every lone pair.
Rule 6 Learn the names of the shapes. The shapes are named from the
position of the atoms and not the position of the orbitals.
CHEMICAL BONDING - LEWIS THEORY
185
TABLE OF SHAPES
Formula
BeCl2
BCl3
CH4
NH3
H2O
Beryllium
chloride
Boron
trichloride
Methane
Ammonia
Water
Bonding Pairs
2
3
4
3
2
Valence Electrons
2
3
4
5
6
Lone Pairs
0
0
0
1
2
Angles between
bonding pairs
180°
120°
109.5°
107°
105°
Name of shape
Linear
Trigonal
Planar
Tetrahedral
Trigonal
Pyramid
Bent
There is one more rule to learn, and it concerns the shape of polyatomic ions.
Rule 2(a) If the molecule is an ion, e.g. ammonium (NH+4), subtract 1 from the number of valence
electrons for every + charge on the ion and add 1 to the valence number for every - charge,
then proceed as before.
SOME MORE EXAMPLES
Formula
NH4+
PCl5
SF6
XeF4
ICI3
Bonding Pairs
4
5
6
4
3
Valence Electrons
5
5
6
8
7
0
0
0
2
2
109.5°
90° & 120°
90°
90°
90°
Tetrahedral
Trigonal
Bipyrimid
Octahedron
Square
T shape
Rule 2(a)
Lone Pairs
Angles between
bonding pairs
Name of shape
5–1=4
186
5 PHYSICAL CHEMISTRY
EXAMINATION QUESTIONS
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
Define or explain the following terms :
(a) Octet rule
(b) Ionic bond
(c) Covalent bond
(d) Co-ordinate covalent bond
(e) Polar covalent bond
(f) Hydrogen bonding
(g) Intermolecular H-bonding
(h) Intramolecular H-bonding
(i) VSEPR theory
(a) Compare the properties of ionic and covalent compounds.
(b) State whether the following compounds are ionic or covalent.
(i) AlCl3
(ii) HgF2
(a) Draw the structure of NaCl crystal and give the co-ordination number of Na+.
(b) Which of the two is more covalent and why in the following pairs
(i) AgCl and AgI
(ii) LiCl and KCl
In methane, ammonia and water molecule the bond angle is decreasing. Explain giving reasons.
(a) Explain the formation of covalent bond between two atoms of chlorine in a chlorine molecule on
the basis of octet rule.
(b) Define (i) Ionic bond; (ii) Co-ordinate bond; and (iii) Metallic bond
(a) What do you understand by ‘Stable configuration’? What are the ways by which an atom can
attain stable configuration?
(b) Write the electronic configuration of any two of the following compounds :
(i) Phosphorus pentachloride
(ii) Sulphuric acid
(iii) Lithium fluoride
What type of bonds do you expect in the following cases? Give reasons :
(i) between a very small cation and a large anion.
(ii) between atoms having a very large difference in electronegativities,
(iii) between atoms of the same element.
Explain qualitatively the valence bond theory with reference to Hydrogen molecule.
Compare the properties of ionic and covalent compounds. Give two examples of each type of compounds.
Indicate the type of bonding that exists in the following solids :
(i) Ice
(ii) Naphthalene
(iii) Diamond
(iv) Potassium chloride
Write Lewis dot formulae of : (a) HOCl (b) BF3 (c) NH4+.
Show the formation of a co-ordinate bond in ozone molecule and discuss briefly the electron gas model
of the metallic bond and how it explains the electrical conductivity of metals.
What is electronegativity? How is the concept of electronegativity used to predict the bond types
between hetero atoms?
Account for the variation of bond angles between the pairs (i) H2O and H2S (104.5° and 92°) and
(ii) H2O and OF2 (104.5° and 101.1°).
Describe the structures of water, ammonia and methane molecules in terms of the electron pair repulsion
theory. Explain why the bond angles are different in the three molecules.
What is a co-ordinate covalent bond? How does it differ from a normal covalent bond?
Discuss the shape of the following molecules on the basis of VSEPR theory :
NH3, CH4, PCl3
Explain the formation of NH3 molecule if no hybridization of s and p-orbitals of nitrogen is assumed.
CHEMICAL BONDING - LEWIS THEORY
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
187
Give diagrammatic representation also.
Explain :
(a) The structure of H2 molecule according to V.B. theory.
(b) Ionic bond and Metallic bond.
Two elements X and Y occur in the same period and their atoms have two and seven valence electrons
respectively. Write down the electronic structure of the most probable compound between X and Y. Will
the bond between X and Y be predominantly ionic or covalent?
Answer. XY2; Ionic
Using VSEPR theory, identify the type of hybridisation and draw the structure of OF2. What are the
oxidation states of O and F?
Account for : The experimentally found N——F bond length in NF3 is greater than the sum of single
covalent radii of N and F.
Which of the following compounds contain bonds that are predominantly ionic in character :
MgO, Ca3P2, AlCl3, Mg2Si and CsF.
Answer. CsF, Mg2Si and MgO
Classify the bonds in the following as ionic, polar covalent or covalent : (a) HCl (b) NaCl and (c) NCl3.
Answer. HCl - Polar covalent, NaCl - Ionic and NCl3 - Covalent
Predict the geometry of the following molecules using VSEPR theory.
(a) CCl4
(b) AlCl3
(c) H2Se
Answer. (a)Tetrahedral (b) Trigonal planar (c) Bent
Predict the geometry of the following ions having VSEPR model.
(b) NO2–
(a) H3O+
–
(c) ClO2
Answer. (a) Pyramidal (b) Bent (c) Bent
Calculate the percentage ionic character of C—Cl bond in CCl4, if the electronegativities of C and Cl are
3.5 and 3.0 respectively.
Answer. 8.875%
The experimentally determined dipole moment, m, of KF is 2.87 × 10–29 coulomb meter. The distance,
d, separating the centers of charge in a KF dipole is 2.66 × 10–3 m. Calculate the percent ionic character
of KF.
Answer. 67.4%
The dipole moment of KCl is 3.336 × 10–29 coulomb meter which indicates that it is highly polar
molecule. The interionic distance between K+ and Cl– in this molecule is 2.6 × 10–10 m. Calculate the
dipole moment of KCl molecule if there were opposite charges of one fundamental unit localised at each
nucleus. Calculate the percentage ionic character of KCl.
Answer. 80%
What is meant by an ionic bond? What are the conditions necessary for the formation of an ionic bond?
(Agra BSc, 2000)
Describe the basic ideas of the VSEPR theory. Explain the application of the theory for predicting the
shapes of the molecules, BCl3, NH3, H2O and SF6.
(Delhi BSc, 2001)
(a) What are electrovalent compounds? Discuss various factors which affect the formation of these
compounds.
(b) What do you understand by hydrogen bonds? Classify them with examples. Explain why water
has abnormally high boiling point.
(Baroda BSc, 2002)
Why bond angles of H2O and NH3 are 104.5° and 107° respectively although central atoms are sp3
hybridized.
(Aligarh BSc, 2002)
Define Lattice energy. Discuss the factors on which it depends.
188
5 PHYSICAL CHEMISTRY
35. (a) Why melting and boiling points of ionic compounds are usually higher than covalent compounds?
(b) Discuss the geometry and shape of PF5 molecule.
(c) Write a short note on hydrogen bonding.
(Arunachal BSc, 2003)
36. Each of the concepts of covalency and electrovalency relates to an idealised state of chemical bonding
which often does not exist in real compounds. Discuss how far this statement is valid and give two
examples with suitable explanation of cases where such non-ideality infact arises. (Delhi BSc, 2003)
37. Strength of hydrogen bond in H–F is more than in H2O but still HF is a gas and H2O is a liquid at room
temperature. Explain.
(Delhi BSc, 2004)
38. (a) The bond angle ∠HNH in ammonia is 107° while bond angle ∠HOH in water is about 104°.
Why?
(b) A covalent bond is stronger than a metallic bond. Why?
(Sambalpur BSc, 2004)
39. Explain intermolecular and intramolecular hydrogen bonding with one example for each.
(Agra BSc, 2005)
40. Based on metallic bond, explain why metals are :
(b) malleable and ductile
(a) good conductors of electricity
(c) having characteristic lustre
(Mysore BSc, 2006)
MULTIPLE CHOICE QUESTIONS
1.
The valency of an element is
(a) the combining capacity of one atom of it
(b) the number of bonds formed by its one atom
(c) the number of hydrogen atoms that combine with one atom of it
(d) all the above
Answer. (d)
2. The octet rule is
(a) the tendency of atoms to have eight electrons in the outermost shell
(b) the tendency of atoms to have eight pairs of electrons in the valency shell
(c) the tendency of the molecule to have a total of eight electrons
(d) the tendency of atoms to have eight non-bonding electrons
Answer. (a)
3. An ionic bond is formed between
(a) two metal atoms
(b) two non-metal atoms
(c) one metal atom and one non-metal atom
(d) one metal atom and one metalloid atom
Answer. (c)
4. Factors governing the formation of an ionic bond are
(a) low ionisation energy of metal and high electron affinity of non-metal atom
(b) high ionisation energy of metal and high electron affinity of non-metal atom
(c) low ionisation energy of metal atom and low electron affinity of non-metal atom
(d) high ionisation energy of metal and low electron affinity of non-metal atom
Answer. (a)
5. The lattice energy is the amount of energy that
(a) is released when one cation combines with one anion
(b) is released when one mole of cations combine with one mole of anions
(c) is released when one mole of an ionic compound is formed from its cations and anions
CHEMICAL BONDING - LEWIS THEORY
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
189
(d) is absorbed when one mole of an ionic compound is formed from its cation and anions
Answer. (c)
The most favourable conditions for the formation of an ionic compound is
(a) low charge on ions, small cation and small anion
(b) high charge on ions, large cation and large anion
(c) high charge on ions, small cation and large anion
(d) low charge on ions, large cation and small anion
Answer. (c)
Ionic compounds are generally
(a) solids having large melting points and good conductors of electricity
(b) gases having low melting points and poor conductors of electricity
(c) solids having low melting points and good conductors of electricity
(d) solids having high melting points and bad conductors of electricity
Answer. (a)
A covalent bond involves
(a) sharing of electrons between a metal and a non-metal atom
(b) sharing of electrons between two metal atoms
(c) sharing of electrons between two atoms having similar electronegativity
(d) sharing of electrons between two atoms having a large difference in electronegativity
Answer. (c)
The total number of electron pairs in a nitrogen molecule is
(a) 2
(b) 3
(c) 5
(d) 7
Answer. (d)
The covalent compounds are soluble in
(a) all acids
(b) all bases
(c) all solvents
(d) non-polar solvents
Answer. (d)
The compounds which contain both ionic and covalent bonds are
(a) CHCl3 and CCl4
(b) KCl and AlCl3
(c) KCN and NaOH
(d) H2 and CH4
Answer. (c)
A co-ordinate bond is formed by
(a) complete transfer of electrons
(b) sharing of electrons contributed by both the atoms
(c) sharing of electrons contributed by one atom only
(d) none of these
Answer. (c)
The types of bonds present in sulphuric acid molecules are
(a) only covalent (b) ionic and covalent
(c) co-ordinate and covalent
(d) co-ordinate, covalent and ionic
Answer. (d)
The common feature among the species O3, SO42–, H3O+ and AlCl3 is that
(a) they contain only ionic bonds
(b) they contain only covalent bonds
(c) they contain co-ordinate bond
(d) they contain covalent and ionic bonds
Answer. (c)
The species CO, CN– and N2 are
(a) isoelectronic
(b) having co-ordinate bond
(c) having low bond energies
(d) having polar bonds
Answer. (a)
190
5 PHYSICAL CHEMISTRY
16. The polarity of a covalent bond is due to
(a) lesser electronegativity difference between two atoms
(b) greater electronegativity difference between two atoms
(c) lesser bond energy
(d) greater bond energy
Answer. (b)
17. A CO2 molecule contains two polar bonds but the net dipole moment is zero. It is because
(a) the molecule has symmetrical linear geometry
(b) the molecule is non-linear
(c) the electronegativity difference between the two atoms is too large
(d) the electronegativity difference between the two atoms is too small
Answer. (a)
18. Among BeF2, BF3, NH3 and CCl4, the molecule with net dipole moment is
(b) BF3
(c) NH3
(d) CCl4
(a) BeF2
Answer. (c)
19. The common feature among the molecules HF, H2O, HCl and NH3 is
(a) intramolecular H-bonding
(b) intermolecular H-bonding
(c) that they contain no polar bonds
(d) that their dipole moment is zero
Answer. (b)
20. Methanol is soluble in water due to
(a) covalent bond nature
(b) ionic bond nature
(c) hydrogen bonding
(d) its poisonous nature
Answer. (c)
21. Among H2O, H2S, H2Se and H2Te, the substance with highest boiling point is
(a) H2O; due to hydrogen bonding
(b) H2S; due to large size of S atom
(c) H2Se; due to large electronegativity difference
(d) H2Te; due to largest size of Te atom
Answer. (a)
22. In ice crystal, the H2O molecules are held together in a
(a) planar structure
(b) linear structure
(c) tetrahedral three dimensional structure
(d) none of these
Answer. (c)
23. The density of ice (solid) is lesser than that of water (liquid) because it has
(a) open cage like structure with no empty spaces
(b) open cage like structure with large empty spaces
(c) intermolecular H-bonding
(d) intramolecular H-bonding
Answer. (b)
24. The density of water is maximum at
(b) 277 K
(c) 281 K
(d) 285 K
(a) 273 K
Answer. (b)
25. Among BeCl2, CHCl3, CCl4 and PCl5, the octet rule is not observed in
(a) BeCl2 only
(b) PCl5 only
(c) BeCl2 and PCl5
(d) CHCl3 and CCl4
Answer. (c)
26. An example of electron deficient compound among BF3, CF4, PF5 and SF6 is
(a) BF3
(b) CF4
(c) PF5
(d) SF6
Answer. (a)
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27. The transition metals show variable valency because of
(a) the availability of vacant d-orbitals
(b) their tendency to form complex ions
(c) their ability to form coloured ions
(d) none of these
Answer. (a)
28. The electrical conductivity of metals is due to
(a) mobile protons is the nucleus
(b) mobile nucleus in the nucleus
(c) mobile electrons in outer vacant spaces
(d) none of these
Answer. (c)
29. According to VSEPR theory,
(a) the lone pairs only decide the structure of the molecule
(b) the bond pairs only decide the structure of the molecule
(c) the lone pairs and bond pairs both decide the structure of the molecule
(d) none of these
Answer. (c)
30. In which of the following, the central atom is surrounded by four electron pairs
(a) H2O
(b) NH3
(c) CH4
(d) All
Answer. (d)
31. The molecule among CCl4, PCl3, SF4 and NH3 that does not contain lone pairs of electrons around the
central atom is
(a) CCl4
(b) PCl3
(c) SF4
(d) NH3
Answer. (a)
32. Which of the following are isostructrual
(a) SO2 and CO2
(b) SO2 and H2O
(c) BCl3 and CHCl3
(d) NH3 and CH4
Answer. (b)
33. The molecular shapes of H2O, NH3 and CH4 are
(a) similar with 2, 1 and 0 lone pairs of electrons respectively
(b) similar with 0, 1 and 2 lone pairs of electrons respectively
(c) different with 0, 1 and 2 lone pairs of electrons respectively
(d) different with 2, 1 and 0 lone pairs of electrons respectively
Answer. (d)
34. The molecule of NH3 is
(a) tetrahedral with bond angle 109° 28′
(b) pyramidal with bond angle 107° 20′
(c) trigonal with bond angle 120°
(d) linear with bond angle 180°
Answer. (b)
35. The NH4+ and SO42– ions have
(a) tetrahedral geometry
(b) triangular geometry
(c) pyramidal geometry
(d) square planar geometry
Answer. (a)
36. Which is incorrect?
(a) all molecules with polar bonds have dipole moment
(b) all molecules with polar bonds may or may not have dipole moment
(c) the greater the difference in electronegativity between two atoms, greater is the polarity
(d) if the electronegativity difference between two atoms is greater than 1.7, the bond will be ionic
Answer. (a)
37. The favourable conditions for the formation of H–bonding are
(a) high electronegativity and small size of the atom bonded to H–atom
(b) low electronegativity and large size of the atom bonded to H–atom
(c) high electronegativity and large size of the atom bonded to H–atom
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5 PHYSICAL CHEMISTRY
(d) low electronegativity and small size of the atom bonded to H–atom
Answer. (a)
The strength of hydrogen bonding lies in between
(a) covalent and ionic bond
(b) metallic and covalent bond
(c) van der Waal’s and covalent bond
(d) metallic and ionic bond
Answer. (d)
The bond angles in a trigonal bipyramid molecules are
(a) 90°
(b) 120°
(c) 109.5°
(d) 120°, 90°
Answer. (d)
CO2 has zero dipole moment whereas H2O has a dipole moment. It is because
(a) H2O is linear while CO2 is a bent molecule
(b) of intermolecular H–bonding in H2O molecules
(c) CO2 is linear while H2O is a bent molecule
(d) CO2 is a gas while H2O is a liquid at room temperature
Answer. (c)
Which of the following does not obey the octet rule?
(a) PCl5
(b) H2O
(c) NH3
(d) CCl4
Answer. (a)
The total number of electrons that take part in forming bonds in O2 is
(a) 2
(b) 4
(c) 6
(d) 8
Answer. (d)
CO is isoelectronic with
(a) C2H2
(b) CN–
(c) O2+
(d) O2–
Answer. (b)
CO2 is isostructural with
(a) H2O
(b) NO2
(c) H2S
(d) C2H2
Answer. (d)
In a bond between two atoms X and Y, the shared electron pair does not lie in the centre. The bond is
(a) single bond
(b) non-polar bond
(c) polar bond
(d) co-ordinate bond
Answer. (c)
The maximum number of Hydrogen bonds formed by a water molecule is
(a) 1
(b) 2
(c) 3
(d) 4
Answer. (b)
Out of the following, intramolecular Hydrogen bonding exists in
(a) water
(b) H2S
(c) 2-nitrophenol
(d) 4-nitrophenol
Answer. (c)
In a compound, hydrogen bonding exists but there is no effect on physical properties like m. pt., b. pt.
etc. It shows the presence of
(a) weak van der Waal’s forces
(b) intramolecular hydrogen bonding
(c) intermolecular hydrogen bonding
(d) resonance in the molecule
Answer. (b)
Which one of the following is the most polar
(a) H — F
(b) H — Cl
(c) H — Br
(d) H — I
Answer. (a)